Chapter 8 – Covalent Bonding

Ms. Wang

Lawndale High School

Chapter 8.1 – Molecular Compounds

In Chapter 7, we learned about electrons being “given up” or “stolen away”

This type of “tug of war” bond between a metal and nonmetal is called an ionic bond

In this chapter, you will learn about another type of bond

Covalent Bond – atoms held together by sharing electrons between two nonmetals

Molecules

We know that a metal cation and nonmetal anion are joined together by an ionic bond and called a SALT

A neutral group of atoms joined together by a covalent bond is called a MOLECULE

Monatomic vs. Diatomic Molecules

Molecules can be monatomic or diatomic

Diatomic Molecule – a molecule consisting of two atoms

There are 7 diatomic molecules on the periodic table (SUPER 7) – N2, O2, F2, Cl2, Br2, I2, H2

Properties of Molecular Compounds

Lower Melting and Boiling Points than Ionic Compounds

Gases or liquids at room temperature

Molecular Formulas

Molecular Formula – the chemical formula of a molecular compound

It shows how many atoms of each element a molecule contains

ExampleH2O contains 3 atoms (2 atoms of H, 1 atom of O)

C2H6 contains 8 atoms (2 atoms of C, 6 atoms of H)

Practice

How many atoms total and of each do the following molecular compounds contain?

1. H2

2. CO

3. CO2

4. NH3

5. C2H6O

Chapter 8.2 – Covalent Bonding

Remember that ionic compounds either gain or lose electrons in order to attain a noble gas electron configuration

Covalent compounds form by sharing electrons to attain a noble gas electron configuration

So the Octet Rule still applies to covalent bonds

Single Covalent Bond

Single Covalent Bond - two atoms held together by sharing one pair of electrons

Unshared Pair / Lone Pair / Nonbonding Pair – a pair of valence electrons that is not shared between atoms

Let’s Practice (you can use dots or lines) F2, H2O, CH4

Double and Triple Covalent Bonds

Double Covalent Bond – a bond that involves two shared pairs of electrons

Triple Covalent Bond – a bond that involves three shared pairs of electrons

Let’s Practice O2

N2

Polyatomic Ion

Polyatomic Ion - tightly bound group of atoms that has a positive or negative charge and behaves as one unit

Some examples are NH4+ and SO3

2-

Bond Dissociation Energy

Bond Dissociation Energy - the energy required to break the bond between two covalently bonded atoms

A large bond dissociation energy corresponds to a strong covalent bond

For example, carbon-carbon has a strong bond dissociation energy so it is not very reactive

Chapter 8.3 - Bonding Theories

Molecular Orbitals – orbitals that apply to the entire molecule instead of just one atom

So far, the orbitals we have been discussing are atomic orbitals (s, p, d, f) for each atom

When two atoms combine, their atomic orbitals overlap and make molecular orbitals

Molecular Orbitals

Just as atomic orbitals belong to a particular atom, a molecular orbital belongs to a molecule as a whole

Each orbital is filled with 2 electrons

Bonding Orbital – a molecular orbital that can be occupied by two electrons of a covalent bond

Sigma Bond ()

Sigma Bond - when 2 atomic orbitals combine to form a molecular orbital that is symmetrical around the axis

S orbitals overlapping P orbitals overlapping end-to-end

Pi Bond ()

Bonding electrons likely to be found in a sausage-shape above and below the axis

Pi bonds tend to be weaker than sigma Pi bonds tend to be weaker than sigma bonds because pi bonds overlap less than bonds because pi bonds overlap less than sigma bondssigma bonds

P orbitals overlapping side-by-side

VSEPR Theory

Explains the 3D shape of molecules

According to VSEPR theory, the repulsion According to VSEPR theory, the repulsion between electron pairs causes molecular between electron pairs causes molecular shapes to adjust so that the valence shapes to adjust so that the valence electron pairs stay as far apart as possibleelectron pairs stay as far apart as possible

A Few VSEPR Shapes

Nine possible molecular shapes

VSEPR Theory

Unshared pairs of electrons (lone pairs) Unshared pairs of electrons (lone pairs) are very important in predicting the shapes are very important in predicting the shapes of moleculesof molecules

ExamplesExamplesMethane (CHMethane (CH44) - tetrahedral) - tetrahedral

Ammonia (NHAmmonia (NH33) - pyramidal) - pyramidal

Water (HWater (H22O) – bentO) – bent

Carbon Dioxide (COCarbon Dioxide (CO22) - linear) - linear

Hybrid Orbitals

VSEPR is good at describing the molecular shapes, but not the types of bonds formed

In hybridization, several atomic orbitals In hybridization, several atomic orbitals mix to form the same total number of mix to form the same total number of hybrid orbitalshybrid orbitals

Orbital hybridization provides information about both molecular bonding and molecular shape

Bond Hybridization

Hybridization Involving Single Bonds – sp3 orbital Ethane (C2H6)

Hybridization Involving Double Bonds – sp2 orbital Ethene (C2H4)

Hybridization Involving Triple Bonds – sp orbital Ethyne (C2H2)

Chapter 7.4 – Polar Bonds and Molecules

There are two types of covalent bonds Polar Covalent Bonds Nonpolar Covalent Bonds

• Polar Covalent Bond – unequal sharing of electrons where one atom has a slightly negative charge and the other atom has a slightly positive charge (HCl, H2O)

• Nonpolar Covalent Bond – equal sharing of electrons between two atoms (Cl2, N2, O2)

Classification of Bonds

You can determine the type of bond artificially by calculating the difference in electronegativity between elements

Type of Bond Electronegativity Difference

Nonpolar Covalent 0 0.4

Polar Covalent 0.5 1.9

Ionic 2.0 4.0

Let’s Practice Together

What type of bond is HCl? (H = 2.1, Cl = 3.1)

Your Turn To PracticeYour Turn To Practice N(3.0) and H(2.1)N(3.0) and H(2.1)

H(2.1) and H(2.1)H(2.1) and H(2.1)

Ca(1.0) and Cl(3.0)Ca(1.0) and Cl(3.0)

Al(1.5) and Cl(3.0)Al(1.5) and Cl(3.0)

Mg(1.2) and O(3.5)Mg(1.2) and O(3.5)

H(2.1) and F(4.0)H(2.1) and F(4.0)

Difference = 3.1 – 2.1 = 1.0Difference = 3.1 – 2.1 = 1.0Therefore it is polar covalent bond.Therefore it is polar covalent bond.

Dipole

• No bond is purely ionic or covalent … they have a little bit of both characters

When there is unequal sharing of electrons a dipole exists Dipole is a molecule Dipole is a molecule that has two poles or that has two poles or regions with opposite regions with opposite chargeschargesRepresented by a Represented by a dipole arrow pointing dipole arrow pointing towards the more towards the more negative end.negative end.

Practice Drawing Dipoles

P- BrP = 2.1Br = 2.8

P –Br P –Br + + --

Practice H(2.1) – S(2.5)H(2.1) – S(2.5) C(2.5) – F(4.0)C(2.5) – F(4.0) Si(1.8) – C(2.5)Si(1.8) – C(2.5) N(3.0) – O(3.5)N(3.0) – O(3.5)

Attractions Between Molecules

dipole interactions – polar molecules attracted dipole interactions – polar molecules attracted to one anotherto one another

dispersion forces – caused by motion of dispersion forces – caused by motion of electrons (weakest of all forces)electrons (weakest of all forces)

Intermolecular attractions are weaker than either ionic or covalent bonds

Van der Waals forces – consists of the two weakest attractions between molecules

Hydrogen Bond

Hydrogen Bonds - attractive forces in Hydrogen Bonds - attractive forces in which a hydrogen covalently bonded to a which a hydrogen covalently bonded to a very electronegative atom is also weakly very electronegative atom is also weakly bonded to an unshared electron pair of bonded to an unshared electron pair of another electronegative atomanother electronegative atom

Hydrogen Bond

This other atom may be in the same This other atom may be in the same molecule or in a nearby molecule, but molecule or in a nearby molecule, but always has to include hydrogenalways has to include hydrogen

Hydrogen Bonds have about 5% of the Hydrogen Bonds have about 5% of the strength of an average covalent bondstrength of an average covalent bond

Hydrogen Bond is the strongest of all Hydrogen Bond is the strongest of all intermolecular forcesintermolecular forces

Intermolecular Attractions

Network Solid – solids in which all of the Network Solid – solids in which all of the atoms are covalently bonded to each otheratoms are covalently bonded to each other

A few solids that consist of molecules do A few solids that consist of molecules do not melt until the temperature reaches not melt until the temperature reaches 10001000ººC or higher called network solids C or higher called network solids (Example: diamond, silicon carbide)(Example: diamond, silicon carbide)

• Melting a network solid would require breaking covalent bonds throughout the solid

Homework

Chapter 8 Assessment Page 247

#’s 39-41, 43-46, 51, 53, 54, 57-59, 61, 65, 68, 83, 85, 86, 89, 96, 99, 100